(And In-Betweens) of Solubility Measurements of Solid Electrolytes*

Pure Appl. Chem., Vol. 85, No. 11, pp. 2077–2087, 2013. http://dx.doi.org/10.1351/PAC-CON-12-11-06 © 2013 IUPAC, Publication date (Web): 20 May 2013 Some highs and lows (and in-betweens) of solubility measurements of solid electrolytes* Glenn Hefter Chemistry Department, Murdoch University, Murdoch, WA 6150, Australia Abstract: Recent solubility measurements of a variety of solid electrolytes in water and aqueous solutions in the author’s laboratories are reviewed. The experimental challenge of per- forming such measurements with high accuracy is demonstrated using the solubility of solid sodium chloride in water at near-ambient temperatures as a paradigm. The special difficulties of measuring low solubilities are demonstrated using Pb(II) sulfate in various aqueous solutions and Pb(II) oxide in sodium hydroxide solutions, and the usefulness of such measurements for obtaining reliable information on homogeneous reactions in solution is briefly discussed. It is also shown, using the alkali metal triflate salts as examples, that determina- tion of the solubilities of even highly soluble salts can be problematic. Lastly, data for the solubilities of a series of sodium carboxylate salts of industrial relevance are discussed and are used to illustrate why the theoretical prediction of solid electrolyte solubilities remains such a challenge. Keywords: aqueous solutions; carboxylate; electrolytes; lead(II); salts; sodium chloride; solubility; triflate. INTRODUCTION The solubilities of substances in solvents are among the oldest of physicochemical measurements [1]. While they would not have been recognized as such at the time, these measurements certainly included the solubilities of solid electrolytes (“salts”) in water. Solubility measurements continue to attract con- siderable attention although the focus nowadays is mostly on systems relevant to practical applications (chemical processing, hydrometallurgy, hydrocarbon extraction, pharmaceuticals, and so on) rather than systematic investigations based on inherent scientific interest [2]. The undiminshed importance of solubility studies is reflected in the continuing success over the last 30 years of the International Symposium on Solubility Phenomena and of the IUPAC-sponsored Solubility Data Series. The latter, which commenced publication in 1979, currently appears under the aegis of the National Institute of Science and Technology (USA) in the Journal of Physical and Chemical Reference Data. The Series will reach its 100th volume within the next year or so, with no end yet in sight. The measurement of the solubility of a simple, stable (chemically unreactive) salt in a stable solvent like water would nowadays be regarded as rather straightforward by most scientists, although it may be noted that the first book-length monograph devoted to the experimental aspects of solubility determinations appeared as recently as 2003 [3], with one of its stated purposes being to improve the quality of the solubility data being reported in the literature. As anyone who has been directly involved in making accurate solubility measurements can attest, there are as many experimental challenges in Pure Appl. Chem. 85, 2027–2144 (2013). A collection of invited papers based on presentations at the 15th International Symposium on Solubility Phenomena and Related Equilibrium Processes (ISSP-15), Xining, China, 22–27 July 2012. 2077 2078 G. HEFTER making reliable solubility determinations as there are in other fields of science. A cursory glance at the literature will show there are often major discrepancies, sometimes at order-of-magnitude levels, among independently determined solubilities, even in relatively uncomplicated systems. The question that then naturally arises from this situation is: how accurately can the solubility of a salt in a solvent be meas- ured, given the simplest possible circumstances of a stable, pure, non-solvate-forming, solid salt of “good” solubility, quantifiable by an inherently reliable technique, and an equally amenable solvent? AN “IN-BETWEEN” SOLUBILITY: NaCl(s) IN H2O Measurement of the solubility of crystalline NaCl in liquid water at ambient conditions of temperature and pressure meets all of the above desiderata. It is an ideal example of an “in-between” solubility that is neither too high nor too low to suffer from the problems associated with such situations (see below). In principle, the solubility of NaCl(s) in H2O(l) should be readily determinable to very high accuracy by evaporative gravimetry, NaCl has a small temperature coefficient of solubility in water, and it does not form hydrates, except at low temperatures [1,4]. More than 100 studies have reported quantitative information on the solubility of NaCl in water at near-ambient conditions [1,4]. A notable result is that determined in the 18th century by Antoine Lavoisier, one of the founders of modern chemistry, which is in respectable agreement with modern values [4]. The reported results have been reviewed on several occasions but the most comprehensive is undoubtedly that of Cohen-Adad and Lorimer [4], published in the IUPAC Solubility Data Series. After a critical assessment of all the original publications: with respect to technique, substance purity, equilibration time, etc., these authors accepted as apparently reliable a total of 41 independently determined results for the solubility of NaCl in H2O at 25 °C. The accepted values showed a spread of 1.4 %, from (26.23 to 26.60) g NaCl/100 g satd. sln. The unweighted average of these values is 26.450 g NaCl/100 g satd. sln., with uncertainties of 1σ = 0.073 (3σ = 0.219) g NaCl/100 g satd. sln. Most physical chemists would be surprised to learn that something as straightforward as the solubility of NaCl(s) in water at 25 °C is known only with a relative certainty of 0.26 % (1σ) or 0.78 % (3σ). But perhaps a better estimate is available? As already noted, there have been other reviews of the solubility of NaCl in water. The numerical values given in those reviews are often encapsulated in the form of an appropriate (usually semi-empirical) mathematical model such as a Pitzer equation. De Visscher and Vanderdeelen [5] have compared the results obtained via these models. They showed (see Fig. 1 in ref. [5]) that at 25 °C the solubility recommended by CODATA/NIST [6] is ~0.75 % higher than the IUPAC value [4], while that given by Archer’s well-known model [7] is ~1.2 % lower. Other models, often parameterized on the same data, fared equally badly; only the value of Pitzer et al. [8] agreed within the (1σ) limit of the experimental uncertainty given above. Thus, it must be concluded that there is not at present a more profound source of NaCl solubility data available than the IUPAC review [4]. But perhaps there is something unusual about the NaCl/H2O system: maybe it is not really a “simple” system after all? To the best of this author’s knowledge, the NaCl/H2O system behaves as “nor- mally” as would be expected, although it does not appear to be widely known that NaCl(s) is rather dif- ficult to dry completely. This somewhat surprising result has been established unequivocally by Rard et al. [9,10] in a series of commendably precise gravimetric experiments. Most pertinent to the present dis- cussion, it was shown that drying NaCl(s) even at a temperature of ~200 °C in a conventional (open) furnace or oven, which encompasses most reported solubility measurements to date (e.g., see below), leaves ~0.1 mass % residual water [9,10]. The amount is higher at lower temperatures. More surpris- ingly, the residual water content is not decreased even with substantial increases in the drying time. It appears that truly dry NaCl(s) can only be obtained (in an open furnace) by heating to 500 °C. Interestingly, Rard et al. reported ongoing weight losses above 550 °C, which they ascribed to surface hydrolysis: © 2013, IUPAC Pure Appl. Chem., Vol. 85, No. 11, pp. 2077–2087, 2013 Solubility measurements of solid electrolytes 2079 NaCl(s) + H2O(g) NaOH(s) + HCl(g) (1) and thereby concluded that there is only a very small temperature range to obtain truly dry, pure NaCl(s), at least in an open furnace. Note, however, that any residual water in NaCl(s) that has been dried to constant weight at temperatures <∼200 °C will only account for a small fraction of the variability observed among independent solubility determinations discussed above. Most of that variability must be ascribed to unknown and unsuspected experimental errors, both random and systematic: a salu- tary lesson on the difficulties of measuring even simple physicochemical quantities with high accuracy. In recent years, the solubility of NaCl(s) in water has been determined routinely in the author’s laboratories as a quality control measure during the study of the solubilities of various salts in water, some of which are described below. This has produced over 80 separate solubility determinations of NaCl(s) in water at various operating temperatures over the last few years. To support the bona fides of these measurements it is appropriate to briefly describe the solubility apparatus used. The form and operation of this equipment has been described in detail elsewhere [11] so only a short summary is given here. The apparatus consists of a transparent thick-walled polycarbonate bath of ~50 L capacity in which is located a plastic carousel, to which up to 28 stoppered, leak-tight polypropylene syringes containing appropriate amounts of solvent and excess solid can be fitted. The carousel is rotated at ~0.1 rpm, causing the excess solid to gently percolate up and down through the solvent without stirring. The bath temperature is controlled to ±0.01 K (short term) but maintains an average temperature to within a few mK over many weeks. A key and distinguishing feature of this apparatus is that it allows the saturated solution to be sampled within the bath, exactly at the equilibration temperature.

Load more